chem 223 exam 1
Which of the following statement regarding a zero order reaction is correct? A. If a reaction is zero order in a reactant, then that reactant is not in the rate determining step of the reaction mechanism. B. If a reaction is zero order in a reactant, then that reactant must be an intermediate. C. If a reaction is zero order in a reactant, then molecules of that reactant are never involved in collisions with molecules of other reactants. D. If a reaction is zero order in a reactant, then that reactant must be a catalyst. E. If a reaction is zero order in a reactant, then the concentration of that reactant does not change during the reaction.
A. If a reaction is zero order in a reactant, then that reactant is not in the rate determining step of the reaction mechanism. zero order reaction: reactant still participates in the reaction, but does not appear in the rate determining step.
Which of the following is the equilibrium constant Kc for the oxidation of copper? 2Cu(s) + O₂(g) ⇌ 2CuO(s) A. Kc = 1/[O₂] B. Kc = [Cu]²[O₂]/[CuO]² C. Kc = [CuO]²/[Cu]²[O₂] D. Kc = 2[CuO]/2[Cu][O₂] E. Kc = [O₂]
A. Kc = 1/[O₂]
A proposed mechanism for the decomposition of NO₂ is: Step 1: NO₂(g) → N(g) + O₂(g) (Slow) Step 2: N(g) + NO₂(g) → N₂O₂(g) (Fast) Step 3: N₂O₂(g) → 2NO(g) (Fast) Which of the following statements is NOT true? A. The activation energy for Step 2 is likely higher than that of Step 1. B. The overall rate law is first order in NO₂. C. The overall reaction is 2NO₂(g) → 2NO(g) + O₂(g). D. The molecularity of the rate-determining step is unimolecular. E. N₂O₂ is an intermediate.
A. The activation energy for Step 2 is likely higher than that of Step 1.
For the equilibrium: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) (ΔH = -197.79 kJ/mol), which of the following disturbances will lead to the consumption of more SO₂? A. The reaction vessel is compressed to half its original volume. B. Some O₂ is removed. More SO₃ is added. C. The temperature increases. D. A catalyst is added. E. Some O₂ is removed.
A. The reaction vessel is compressed to half its original volume.
For which of the following reactions are the values of Kc and Kp equal? (Multiple choice - choose all right answers!) A. 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) B. H₂O(g) + CO(g) ⇌ H₂(g) + CO₂(g) C. C(s) + O₂(g) ⇌ CO₂(g) D. H₂O(l) + CO(g) ⇌ H₂(g) + CO₂(g)
B. H₂O(g) + CO(g) ⇌ H₂(g) + CO₂(g) C. C(s) + O₂(g) ⇌ CO₂(g)
A three-step mechanism has been proposed for a reaction: Step 1: Cl₂ ⇌ 2Cl (fast and reversible) Step 2: Cl + CHCl₃ → HCl + CCl₃ (slow) Step 3: Cl + CCl₃ → CCl₄ (fast) Which of the following statements is INCORRECT? A. The overall reaction is first order in CHCl₃. B. The rate-determining step is unimolecular. C. The overall reaction is: Cl₂ + CHCl₃ → HCl + CCl₄ D. The activation energy of Step 2 is likely higher than that of Step 1. E. Cl is an intermediate.
B. The rate-determining step is unimolecular.
For the equilibrium: CH₄(g) + H₂O(g) ⇌ CO(g) + 3H₂(g) (ΔH = 206 kJ/mol), which of the following disturbances will lead to the formation of more CH₄? A. A catalyst is added. B. The temperature decreases. C. More H₂O is added. D. Some CO is removed. E. The reaction vessel expands to three times its original volume.
B. The temperature decreases.
Two reactions, each of which only has one single reactant, have the same rate constant. The first reaction is first order, and the second reaction is second order. If the initial concentrations of the reactants are both 2.0 M, which reaction will proceed at the higher rate? A. the first reaction B. the second reaction C. cannot be determined D. their rates are the same
B. the second reaction rate = k x [reactant]^order. In this case: k is the same for the two reactions. rate of reaction 1: k x (2.0M), rate of reaction 2: k x (2.0M)² = k x 4.0M², so rate 2 > rate 1
The exothermic reaction 2NO(g) ⇌ N₂(g) + O₂(g) is at equilibrium at 1400 K, with an equilibrium constant 2.6 × 10⁵. Which of the following statement is CORRECT? A. The reaction proceeds mostly in the reverse direction. At equilibrium, there is almost no products left. B. The equilibrium constant of the reaction increases as the temperature increases. C. kf is greater than kr. D. The forward reaction rate is greater than the reverse reaction rate at equilibrium. E. At equilibrium, there is no net reaction, so both the forward and the reverse reaction have stopped.
C. kf is greater than kr.
For the equilibrium: 2NOBr(g) ⇌ 2NO(g) + Br₂(g), Kc = 1.56 × 10⁻³ at 300 K. In a reaction vessel, the concentrations are [NOBr] = 0.3732 M, [Br₂] = 0.0134 M, and [NO] = 0.0268 M. In which direction will the equilibrium proceed? A. To the left B. To make more NOBr C. Cannot be determined D. To make more Br₂ E. The reaction is already at equilibrium
D. To make more Br₂
Which of the following statements about catalysts is CORRECT? A. A homogeneous catalyst appears in the overall rate law of the reaction. B. A catalyst speeds up the reaction by increasing the activation energy. C. A catalyst is the same as an intermediate. D. A catalyst that is in a different phase from the reactants is a homogeneous catalyst. E. A catalyst may appear in some of the elementary steps of a reaction.
E. A catalyst may appear in some of the elementary steps of a reaction. A catalyst speeds up reaction by decreasing the activation energy. It may appear in some elementary steps, but should not appear in the overall reaction. It is different from an intermediate, because an intermediate does not speed up the reaction.
T/F - As the reaction goes on, the instantaneous rate of the reaction increases.
False
(a) For the equilibrium: N₂O₄(g) ⇌ 2NO₂(g), K₁ = 1.5 × 10⁶. What is the equilibrium constant for 2NO₂(g) ⇌ N₂O₄(g)? (b) For the equilibrium: ½N₂(g) + O₂(g) ⇌ NO₂(g), K₂ = 1.2 × 10⁻⁵. What is the equilibrium constant for N₂(g) + 2O₂(g) ⇌ 2NO₂(g)? (c) Based on your answers in (a) and (b), what is the equilibrium constant for N₂(g) + 2O₂(g) ⇌ N₂O₄(g)?
K₂ = 6.67 x 10⁻⁷ K₁ = 1.44 x 10⁻¹⁰ K = 9.6 x 10⁻¹⁷
For the equilibrium 2NOBr(g) ⇌ 2NO(g) + Br₂(g), Kc = 1.56 × 10⁻³ at 300 Kc. Suppose 0.400 M NOBr was initially placed in a reaction vessel at 300 K. At a given point, the concentration of Br₂ is 0.0134 M. Calculate the reaction quotient, and determine whether the reaction is at equilibrium.
No, the equation is not at equilibrium because it proceeds in the forward direction
The equilibrium constant Kc for the hypothetical reaction 2C ⇌ D + E is 3 × 10⁻³. At a particular time, the composition of the reaction mixture is [C] = 0.02 M and [D] = [E] = 0.01 M. In which direction will the reaction proceed?
Reverse
Given the following data for the reaction: 2NO(g) + Cl₂(g) → 2NOCl(g): Experiment 1 [NO] 0.0400 (M), [Cl2] 0.0150 (M), Rate 3.4 × 10⁻⁴ (M/s) Experiment 2 [NO] 0.0100 (M), [Cl2] 0.0150 (M), Rate 8.5 × 10⁻⁵ (M/s) Experiment 3 [NO] 0.0100 (M), [Cl2] 0.0300 (M), Rate 3.4 × 10⁻⁴ (M/s) Determine: (a) The rate law. (b) The rate constant.
a) b)
A proposed mechanism for the reduction of nitrogen as NO by hydrogen is: Step 1: H₂(g) + 2NO(g) → N₂O(g) + H₂O(g) (Slow) Step 2: N₂O(g) + H₂(g) → N₂(g) + H₂O(g) (Fast) (a) Sketch the energy vs. progress of reaction from the reactants to the products of the overall reaction. A rough sketch by hand is fine, however, make your own sketch! Also, make sure you label the axes, and show the relative heights of the activation energies of the two steps. (b) On your sketch in (a), mark the activation energy for each step, the activated complex of each step, the intermediate, and the energy of the overall reaction. (c) On your sketch in (a), using dashed lines, sketch what the energy vs. progress of reaction would look like if a catalyst is added to the reaction.
a) b) c)
For the equilibrium H₂O(g) + CH₄(g) ⇌ CO(g) + 3H₂(g), when 1.00 atm CH₄ and 1.00 atm H₂O are placed in a reaction vessel at 1050 K, their equilibrium pressures are both 0.530 atm. (a) Calculate Kp at 1050 K. (b) Calculate Kc at 1050 K. (c) Given that ΔHrxn = 206.1 kJ/mol, calculate Kp at 298 K. (d) Calculate ΔGrxn° at 298 K.
a) b) c) d)
For the equilibrium reaction: CH₄(g) + 2H₂S(g) ⇌ CS₂(g) + 4H₂(g), the concentration at equilibrium are [CH₄] = 0.3322 M, [H₂S] = 0.6644 M, [CS₂] = 0.0678 M, and [H₂] = 0.2712 M at 1400 K. (a) Calculate Kc at 1400 K. (b) Calculate Kp at 1400 K. (c) If ΔGrxn⁰ = 181.92 kJ/mol at 298 K, calculate Kp at 298 K. (d) Assuming that ΔHrxn⁰ stays constant from 298 K to 1400 K, calculate ΔHrxn⁰.
a) b) c) d)
For the reaction: 4NH₃(g) + 7O₂(g) → 4NO₂(g) + 6H2(g), if Δ[H₂]/Δt = 0.0035 M/s, what is the rate of consumption of O₂?
-0.0041 M/s
For the reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), Kp = 4.3 × 10⁻³ at 300 °C and 4.5 × 10⁻⁵ at 450 °C. What is the value of ΔHrxn°?
-104.75 kJ/mol ln(k1/k2) → ÷ (1/T₂-1/T₁) → x R → ÷ 1000
For the equilibrium: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), at 830 °C, A vessel that initially contains a mixture of SO₂ and O₂, with the partial pressure of SO₂ being 1.25 atm and the partial pressure of O₂ being 0.63 atm, was allowed to come to equilibrium. If the partial pressure of SO₃ at equilibrium is 0.10 atm, what is the equilibrium constant of the reaction at 830 °C?
0.013
The value of Kc for the reaction: CO(g) + H₂O(g) ⇌ CO₂(g) + H₂(g) at 700 K is 5.1. If a mixture initially has both the concentration of CO and the concentration of H₂O being 0.050 M, what is the equilibrium concentration of CO₂?
0.035 M
For the equilibrium: N₂(g) + O₂(g) ⇌ 2NO, Kc = 0.050 at 2200 °C. If a reaction vessel initially contained 0.21 atm N₂ and 0.21 atm O₂, calculate the equilibrium partial pressure of NO.
0.0422 atm
At 300 °C, the equilibrium constant Kc for the reaction: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) is 2.4 × 10⁻³. What is the value of the equilibrium constant for SO₂(g) + ½O₂(g) ⇌ SO₃(g)?
0.049
For the reaction 2O₃(g) → 3O₂(g), if the concentration of O₂ increased from 0.0250 M to 4.15 M over the first 5.50 s, what is the average rate of the reaction during this time?
0.25 M/s
The reaction: 2NOCl(g) ⇌ 2NO(g) + Cl₂(g) has a Kp value of 1.6 × 10⁻⁵ at some temperature. If 0.500 atm of NOCl is placed in a closed vessel and allowed to come to equilibrium, what is its equilibrium partial pressure? (Make the "x is small" assumption)
0.48 atm
For the reaction 2NO₂(g) → 2NO(g) + O₂(g), the plot of 1/[NO₂] vs. t is linear. If the initial concentration of NO₂ is 0.500 M, and the half-life of the reaction is 3.67 s, what is the rate constant of the reaction?
0.545 M⁻¹ s⁻¹
Given the following data for the reaction: CO(g) + Cl₂(g) → COCl₂(g): Experiment 1: [CO] 0.10 (M), [Cl₂] 0.20 (M), Rate 0.1008 (M /s) Experiment 2: [CO] 0.20 (M), [Cl₂] 0.20 (M), Rate 0.2016 (M /s) Experiment 3: [CO] 0.20 (M), [Cl₂] 0.60 (M), Rate 1.048 (M /s) What is the order of the reaction with respect to CO?
1 Experiments 1 and 2 should be used to calculate the order of CO. log(0.2016/0.1008)/log(0.20/0.10) = 1 The reaction is 1st order in CO.
For the reaction: 2ClO(g) → Cl₂(g) + O₂(g)If Δ[ClO]/Δt at 298 K is -2.3×10⁷ M/s, what is Δ[Cl₂]/Δt?
1.15×10⁷ M/s
The reaction: 2NO + 2H₂ → N₂ + 2H₂O is first order in H₂ and second order in NO. At a certain temperature, when the initial concentration of NO is 0.20 M, and the initial concentration of H₂ is 0.10 M, the initial rate is 0.00492 M s⁻¹. What is the rate constant at that temperature?
1.23 M⁻² s⁻¹ rate = k * [H₂] * [NO]² Plug in the numbers: rate = 0.00492 M s⁻¹ = k * (0.10 M) * (0.20 M)² k = 1.23 M⁻² s⁻¹
For the equilibrium: 2ClO(g) ⇌ Cl₂O₂(g), Kc = 4.96 × 10¹¹ at 253 K. What is Kc at 253 K for the equilibrium: ½ Cl₂O₂(g) ⇌ ClO(g)?
1.42 × 10⁻⁶
Given the following data for the reaction: CO(g) + Cl₂(g) → COCl₂(g): Experiment 1: [CO] 0.10 (M), [Cl₂] 0.20 (M), Rate 0.1008 (M /s) Experiment 2: [CO] 0.20 (M), [Cl₂] 0.20 (M), Rate 0.2016 (M /s) Experiment 3: [CO] 0.20 (M), [Cl₂] 0.60 (M), Rate 1.048 (M /s) What is the order of the reaction with respect to Cl₂?
1.5 Experiments 2 and 3 should be used to calculate the order of Cl₂. log(1.048/0.2016)/log(0.60/0.20) = 1.5 The reaction is 1.5 order in Cl₂. The correct answer is: 1.5
For the equilibrium reaction: 2NO(g) + O₂(g) ⇌ 2NO₂(g), if ΔGrxn° = -69.7 kJ/mol at 298 K, calculate Kp at 298 K.
1.65 × 10¹²
For the reaction: CaO(s) + SO₂(g) ⇌ CaSO₃(s), Kp = 5.2 × 10³ at 300 K. If all gas pressures are measured in the unit of atm, what is the equilibrium pressure of SO₂?
1.92 × 10⁻⁴ atm
The decomposition reaction of NO₂: 2NO₂(g) → 2NO(g) + O₂(g) follows the rate law: rate = k[NO₂]². Which of the following plot is linear? A. ln(1/[NO₂]) vs. t B. [NO₂] vs. t C. 1/[NO₂] vs. t D. [NO₂]² vs. t E. ln[NO₂] vs. t
1/[NO₂] vs. t Second order reaction: the plot of 1/[reactant] vs. t is linear.
Most disinfecting wipes claim they kill 99.99% of the germs. If it takes 10 seconds to kill half of the germs on a counter top after wiping it with a disinfecting wipe, how long will it take to kill 99.99% of the germs?
132.9 s
For the reaction: 2INO(g) → I₂(g) + 2NO(g), if Δ[INO]/Δt = -2.2 × 10⁴ M/s, calculate the rate of formation of I₂ and the rate of formation of NO.
2.2 × 10⁴ M/s
The rate constant for the first-order decomposition of a particular hydrocarbon is 1.85 × 10⁻⁶ s⁻¹ at 485 K. If the activation energy is 84.0 kJ/mol, what is the frequency factor at this temperature?
2062.2 s⁻¹
Azomethane (CH₃NNCH₃) is a heat-sensitive explosive that decomposes to yield nitrogen gas and methyl radicals: CH₃NNCH₃(g) → N₂(g) + 2CH₃(g). At 600 °C, the reaction has an activation energy of 21.4 kJ/mol and k = 2.00 × 10¹⁸ s⁻¹. What is the frequency factor, A, for this reaction?
3.81 × 10¹⁹ s⁻¹ k = Ae^(-Ea/RT), plug in the numbers given for k, Ea, and T, and solve for A
For the rate law: rate = k[A]³'²[B], the order of the reaction in A is _____, the order of the reaction in B is _____, and the total order is _____.
3/2, 1, 5/2
For the reaction: 2NO(g) + O₂(g) ⇌ 2NO₂(g), at equilibrium, the concentrations of each species in a sealed reaction vessel are [NO] = 0.004 M, [O₂] = 0.007 M, and [NO₂] = 0.006 M. Calculate Kc of the reaction.
321.4
Nitrous acid slowly decomposes to NO, NO₂, and water in the following reaction: 2HNO₂(aq) → NO(g) + NO₂(g) + H₂O(l). Use the data below to determine the average rate of the reaction between 1500 and 2000 minutes: Time 0 (min); [HNO₂] 0.1560 (μM) Time 1000 (min); [HNO₂] 0.1466 (μM) Time 1500 (min); [HNO₂] 0.1424 (μM) Time 2000 (min); [HNO₂] 0.1383 (μM)
4.1 × 10-⁶ μM/min Δ[HNO₂]/Δt = (0.1383 - 0.1424)/(2000-1500) = -8.2 × 10-⁶ μM/min rate of the reaction = (-1/2) Δ[HNO₂]/Δt = 4.1 × 10-⁶μM/min
The rate law of a hypothetical reaction: A + B + C -> D is: Rate = k[A][B]³[C]⁰∙⁵ . What is the overall order of the reaction?
4.5 overall order = 1+3+1/2 = 4.5
For the decomposition of N₂O₅, the rate constant is 2.14 × 10⁵ s⁻¹ at 658 K, and the activation energy is 101.608 kJ/mol. What is the rate constant at 688 K?
4.81 × 10⁵ s⁻¹
If the value of Kp for the reaction: Cl₂(g) + CO(g) ⇌ COCl₂(g) is 0.10 at 327 °C, what is the value of Kc at 327 °C? (Use R = 0.08206 L.atm/mol.K)
4.9
At 300 °C, the equilibrium constant Kc for the reaction: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) is 2.4 × 10⁻³. What is the value of Kc for 2SO₃(g) ⇌ 2SO₂(g) + O₂(g)?
417
For the reaction: NO(g) + F₂(g) → NOF(g) + F(g), the rate constant is 4.72 × 10⁷ M⁻¹s⁻¹ at 298 K. If the activation energy is 6.31 kJ/mol, at what temperature will the rate constant be 1.21 × 10⁸ M⁻¹s⁻¹?
472.7 K
The decomposition reaction of NO₂: 2NO₂(g) → 2NO(g) + O₂(g) follows the rate law: rate = k[NO₂]². If the rate constant is 0.544 M⁻¹ s⁻¹ at a particular temperature and the initial concentration of NO₂ is 0.0100 M, how many seconds will it take for the concentration of NO₂ to reach 6.25% of its original concentration?
735 s
The equilibrium constant Kc for the reaction: N₂O₄(g) ⇌ 2NO₂(g) is 4.1 × 10⁻⁵ at 0 °C. If a reaction vessel at this temperature initially contains 0.022 M N₂O₄, what will be the partial pressure of NO₂ in the vessel when the reaction reaches equilibrium? (Make the "x is small" assumption)
9.50 × 10⁻⁴ M
Each of the following reactions is first order with respect to each reactant. Which reaction has the fastest initial rate if the initial concentrations of the reactants are the same? A. ClO(g) + NO(g) → NO₂(g) + Cl(g), k = 1.7 × 10⁻¹¹ L/(mol.s) B. ClO₂(g) + O₃(g) → ClO₃(g) + O₂(g), k = 3.0 × 10⁻¹⁹ L/(mol.s) C. ClO(g) + O₃(g) → ClO₂(g) + O₂(g), k = 1.5 × 10⁻¹⁷ L/(mol.s) D. ClO₂(g) + NO(g) → NO₂(g) + ClO(g), k = 3.4 × 10⁻¹³ L/(mol.s)
A. ClO(g) + NO(g) → NO₂(g) + Cl(g), k = 1.7 × 10⁻¹¹ L/(mol.s) Rate = k x [reactant]^order. If concentrations and reaction orders are the same, then the reaction with the largest k has the highest rate.
Which of the following reactions, when at equilibrium, will have proceeded to the right the most at the given temperature? A. H₂(g) + Cl₂(g) ⇌ 2HCl(g), Kp = 3.8 × 10⁴ at 1000 K B. H₂(g) + I₂(g) ⇌ 2HI(g), Kp = 0.034 at 1000 K C. P₂(g) + 3H₂(g) ⇌ 2PH3(g), Kp = 12.89 at 873 K D. CO(g) + 3H₂(g) ⇌ CH₄(g) + H₂O(g), Kp = 1.61 × 10⁻⁵ at 1400 K E. 2NOCl ⇌ Cl₂(g) + 2NO(g), Kp = 0.0193 at 500 K
A. H₂(g) + Cl₂(g) ⇌ 2HCl(g), Kp = 3.8 × 10⁴ at 1000 K
Nitrous oxide (N₂O) is used as an anesthetic (laughing gas) and in aerosol cans to produce whipped cream. It is a potent greenhouse gas and decomposes slowly into N₂ and O₂: 2N₂O(g) → 2N₂(g) + O2(g) (a) If a plot of ln[N₂O] vs. time is linear, what is the order of the reaction with respect to N₂O? (b) If the rate constant of the reaction is 6.8 × 10⁻³ s⁻¹ , what is the half-life of the reaction? (c) How many seconds will it take for the concentration of N₂O to reach 1.56% of its original concentration?
a) 1st b) 101.95 c) 611.65
Which of the following statements about chemical equilibrium is CORRECT? Select one: a. The reactants have converted 100% to products when the reaction reaches equilibrium. b. The equilibrium constant of an exothermic reaction increases as the temperature increases. c. At equilibrium, the concentrations of the reactants and products are equal. d. At equilibrium, the rate of the forward reaction and the rate of the reverse reaction are equal. e. A reaction stops when it reaches equilibrium.
d. At equilibrium, the rate of the forward reaction and the rate of the reverse reaction are equal.
NO₂(g) + CO(g) → NO(g) + CO₂(g) Experiment 1 [NO₂] 0.263 (M), [CO] 0.826 (M), Initial rate 1.44 × 10⁻⁵ (M /s) Experiment 2 [NO₂] 0.263 (M), [CO] 0.413 (M), Initial rate 1.44 × 10⁻⁵ (M /s) Experiment 3 [NO₂] 0.526 (M), [CO] 0.413 (M), Initial rate 5.76 × 10⁻⁵ (M /s) a) Determine the rate law. b) Determine the order with respect to each reactant, and the overall reaction order. c) Determine the rate constant.
k[NO₂]² 2nd order 2.082 x 10⁻⁴ M⁻¹/s
